Le Chatelier's principle and industrial applications

Chemistry · Chemical equilibrium. The idea in plain words, the rules to remember, the mistake to avoid and a worked example, written and checked by Mega Lecture.

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The idea

Le Chatelier's principle says: if a system at equilibrium is disturbed, it behaves in such a way as to cancel the effect of that disturbance. So the equilibrium moves in the direction that undoes the change you made.

Adding a reactant moves it to the right. Raising the pressure moves it to the side with fewer gas molecules. Raising the temperature moves it in the endothermic direction. Only a change of temperature changes the value of K_c.

A catalyst speeds up the forward and reverse reactions equally. Equilibrium is reached sooner, but its position and K_c stay the same. Industry chooses conditions that balance yield, rate and cost.

Rules to remember

Concentration: adding a reactant (or removing a product) shifts the equilibrium to the right; K_c does not change.
Pressure: higher pressure shifts the equilibrium to the side with fewer moles of gas; no effect when the moles of gas are equal on both sides.
Temperature: heating favours the endothermic direction, cooling favours the exothermic direction.
Exothermic forward reaction: K_c decreases as temperature rises. Endothermic forward reaction: K_c increases.
Haber process: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ mol⁻¹; about 200 atm, 400–450 °C, finely divided iron catalyst; ammonia is removed by liquefying it.
Contact process: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), exothermic; 400–500 °C, V₂O₅ catalyst, pressure only slightly above atmospheric.
The temperature used in both processes is a compromise: low temperature gives a better yield but too slow a rate.

Common mistake

Students say that a change of pressure or concentration changes K_c. It does not: the position of equilibrium moves, but only temperature changes the value of K_c.

Worked example

N₂O₄(g, colourless) ⇌ 2NO₂(g, brown) is endothermic in the forward direction. The mixture is cooled at constant pressure. What happens to the colour and to K_c?

Cooling removes heat, so the equilibrium moves in the direction that gives out heat (the exothermic direction).
The forward reaction is endothermic, so the reverse reaction is exothermic: the equilibrium shifts to the left.
NO₂ is used up and N₂O₄ forms, so the brown colour becomes paler.
K_c = [NO₂]² ÷ [N₂O₄]; less product and more reactant means a smaller value.

Answer: The brown colour fades and K_c decreases.

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Chapter 8: Chemical Equilibrium

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